Electrochemical Synthesis of Ammonia in Solid Electrolyte Cells
Ioannis Garagounis, Vasileios Kyriakou, Aglaia Skodra, Eirini Vasileiou and Michael Stoukides
Journal Name:
Frontiers in Energy Research
ISSN:
2296-598X
Article type:
Review Article
Received on:
25 Nov 2013
Accepted on:
04 Jan 2014
Provisional PDF published on:
04 Jan 2014
Frontiers website link:
www.frontiersin.org
Citation:
Garagounis I, Kyriakou V, Skodra A, Vasileiou E and Stoukides
M(2014) Electrochemical Synthesis of Ammonia in Solid Electrolyte
Cells. 2:1. doi:10.3389/fenrg.2014.00001
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name=fuel%20cells&ART_DOI=10.3389/fenrg.2014.00001
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Frontiers in Fuel Cells
25th November 2013
Electrochemical Synthesis of Ammonia in Solid Electrolyte Cells
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I. Garagounis1,2, V. Kyriakou1,2, A. Skodra2, E. Vasileiou1,2, M. Stoukides1,2*
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Department of Chemical Engineering, Aristotle University of Thessaloniki
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2
Chemical Processes and Energy Resources Institute, CERTH
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*Correspondence: Michael Stoukides, Aristotle University of Thessaloniki, Department of
Chemical Engineering, Laboratory of Electrochemical Processes, Thessaloniki, 54124, Greece.
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stoukidi@cperi.certh.gr
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Keywords: ammonia synthesis, solid electrolytes, proton conductors, oxygen conductors,
solid state ammonia synthesis
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Abstract
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Developed in the early 1900's, the “Haber-Bosch” synthesis is the dominant NH3 synthesis process.
Parallel to catalyst optimization, current research efforts are also focused on the investigation of
new methods for ammonia synthesis, including the electrochemical synthesis with the use of solid
electrolyte cells. Since the first report on Solid State Ammonia Synthesis (SSAS), more than 30
solid electrolyte materials were tested and at least 15 catalysts were used as working electrodes.
Thus far, the highest rate of ammonia formation reported is 1.13×10−8 mol s−1 cm−2, obtained at
80°C with a Nafion solid electrolyte and a mixed oxide, SmFe0.7Cu0.1Ni0.2O3, cathode. At high
temperatures (>500oC) the maximum rate was 9.5*10-9 mol s−1 cm−2 using Ce0.8Y0.2O2-δ -[Ca3(PO4)2
-K3PO4] as electrolyte and Ag-Pd as cathode. In this paper, the advantages and the disadvantages
of SSAS vs the conventional process and the requirements that must be met in order to promote the
electrochemical process into an industrial level, are discussed.
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1. Introduction
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In September 2013, one hundred years were completed since the first industrial reactor for
ammonia synthesis was put in operation. Designed by Carl Bosch and still standing in a garden in
front of the BASF Ammonia Synthesis Institute, this high-pressure reactor could produce five tons
of ammonia per day [1]. With the fertilizer industry being by far its primary user, currently, the
worldwide production of ammonia is over one hundred million tons per year [2].
33
The development of an industrial process for ammonia synthesis from its elements:
34
N2 + 3 Η2 <==> 2 NH3
(1)
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
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is considered the bellwether reaction of heterogeneous catalysis [3] and has played a central role in
the development of the chemical industry during the 20th century [4]. The Haber-Bosch process,
which involves reaction of gaseous nitrogen and hydrogen on a Fe-based catalyst at high pressures
(150 - 300 bar) was developed in the early 1900's after an extensive search for materials with an
industrially acceptable catalytic activity [4, 5]. The nitrogen molecule, N2, is one of the most stable
molecules in nature; the triple bond between the two N atoms has a binding energy of 962 kJ/mol.
Moreover, in addition to the difficulty in dissociating N2, the reaction conversion is limited by
thermodynamics. The gas volume decreases with reaction. Therefore, very high pressures must be
used in order to push equilibrium to the right according to the Le Chatelier principle. In the early
stages of the development of the process, very high pressures (500-1000 bar) were used [5]. The
equilibrium conversion decreases with temperature because the reaction is exothermic (-46 kJ/mol
at 298 K). Hence, if thermodynamics only were taken into account, it would be preferable to operate
at room temperature at which the predicted conversion would be higher than 90% at atmospheric
pressure. The reaction kinetics, however, are unacceptably slow at room temperature. Thus, the
trade-off solution was to operate at temperatures near 450oC and at pressures between 150-300 bar
in order to obtain an equilibrium conversion of the order of 10 to 15% [2]. It is beyond the purpose
of this communication to review the research that has been conducted thus far concerning the
kinetics and the mechanism of ammonia synthesis. Numerous papers can be found in literature.
Fortunately, earlier and recent works have been reviewed in the last three decades [1, 4].
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Nature converts molecular N2 into NH3 without the use of high temperatures and pressures
[6-8]. This natural process utilizes metalloenzymes and has been developed and optimized over a
period of a few billion years [6]. The enzymes, called nitrogenases, contain iron and molybdenum
and catalyze the reaction of ammonia synthesis from atmospheric nitrogen, protons and electrons
[6, 8]. The high efficiency of this enzyme-catalyzed synthesis has attracted the interest of
researchers in imitating nature. In the last two decades, remarkable progress has been achieved in
both, understanding the mechanism of the biological process and developing an “artificial”
abiological process using transition metal complexes [8-10].
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Since its discovery, the Haber-Bosch process has been substantially improved and
optimized. The continuous search for more active catalysts resulted in operation at lower
temperatures and consequently, at lower pressures. Parallel to catalyst optimization, research
efforts also focus on the investigation of new methods for ammonia synthesis, including the
electrochemical synthesis in aqueous or solid electrolyte cells [11]. Several studies in which
ammonia was electrochemically produced in aqueous media have been reported [12, 13]. One
problem, however, with aqueous electrolytes is that the operating temperature must be low. At low
temperature, the kinetic reaction rates are inevitably slow. The discovery and development of solid
state proton (H+) conductors [14] offered the opportunity to produce ammonia at elevated
temperatures. The first work on Solid State Ammonia Synthesis (SSAS) was reported in 1998 [15].
Since then, groups worldwide have studied this subject and a number of problems that should be
solved in order to enable scale-up, have been identified. A comprehensive review of the SSAS
studies reported until the end of 2010 was published by Amar et al. [2]. The cell component
requirements were discussed with emphasis on the solid electrolyte materials. Recently, Giddey et
al. [16] carefully reviewed the progress on the electrochemical production of ammonia with
emphasis on materials of construction, major technical challenges and status of technology.
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The present review focuses primarily on SSAS. The properties and requirements to be met
by the working electrode, which also serves as the ammonia synthesis catalyst, are discussed in
more detail. The electrochemical results are compared to those obtained by the catalytic process.
Finally, the hurdles to be overcome in order to bring SSAS into industrial practice are discussed.
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Garagounisetal.
ElectrochemicalSynthesisofAmmonia
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2. Ammonia Synthesis in Proton (H+) Conducting Cells
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Figure 1 shows schematically the principle of SSAS from gaseous N2 and H2 in a protonconducting solid electrolyte cell. The cell consists of a solid state H+ conductor. Two porous metal
films are placed on the two sides of the solid electrolyte and serve as electrodes. Gaseous Η2 passes
over the anode where it is converted to protons (H+) according to the reaction:
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3 Η2 <==> 6 H+ + 6 e-
(2).
In the form of protons, hydrogen is transported through the electrolyte to the cathode, where the
half-cell reaction:
N2 + 6 H+ + 6 e-
<==>
2 NH3
(3)
takes place. Thus, reaction (1) is again the overall reaction with the cathodic electrode serving as
catalyst for ammonia synthesis.
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Figure 1. Schematic diagram of a Solid State H+ Conducting
Cell Used for NH3 Synthesis from its Elements.
The use of proton conducting solid electrolyte cells to bypass the thermodynamic
constraints of reversible hydro- or dehydrogenation reactions was proposed in the mid 1990s [1719]. Panagos et al. [17] developed a model which showed that, if a H+ cell is used, the product yield
of a hydrogenation reaction can exceed that of the corresponding catalytic reactor by several orders
of magnitude. Marnellos et al. used that model to examine the performance of various types of
electrochemical reactor design and discussed the parameters that affect the reaction conversion [1821].
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
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Experimental verification of SSAS was first demonstrated in 1998 [15]. The solid
electrolyte cell used was of the form: H2/Pd/SCY/Pd/N2. The H+ conductor (SCY) was a strontiaceria-ytterbia perovskite of the form SrCe0.95Yb0.05O3. Two porous polycrystalline Pd films were
used as electrodes. At 570oC and atmospheric pressure, 78% of the supplied protons were
converted into ammonia.
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In the past fifteen years, a large number of researchers have studied this system. Table 1
contains the SSAS works published until September 2013. The first column on the left-hand side
shows the formula (and its abbreviated name in parentheses) of the solid electrolyte that was used.
Similarly, the second column shows the working electrode (WE) used. The third column contains
the reference of each work. In all studies of Table 1, reactions (2) and (3) occurred at the anode and
cathode, respectively. Hence, ammonia was formed at the cathode (WE).
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Table 1: The proton conducting electrolytes and electrodes used by various investigators in solid
state ammonia synthesis cells
Solid Electrolyte
Working Electrode (WE)
Reference
SrCe0.95Yb0.05O3-a (SCY)
Pd
15,21
CaZr0.90In0.1O3-a (CZI)
Fe
22
SrZr0.90Y0.1O3-a (SZY)
Fe
23
Ba3(Ca1.18Nd1.82)O9-δ (BCN18)
Ag-Pd
24
Ba3CaZr0.5Nd1.5O9-δ (BCZN)
Ag-Pd
24
Ba3Ca0.9Nd0.28Nb1.82O9-δ (BCNN)
Ag-Pd
24
BaCe0.9Sm0.1O3-a (BCS)
Ag-Pd
25
La0.9Sr0.1Ga0.8Mg0.2O3-δ (LSGM)
Ag-Pd
26
La0.9Ca0.1Ga0.8Mg0.2O3-δ (LCGM)
Ag-Pd
27
La0.9Ba0.1Ga0.8Mg0.2O3-δ (LBGM)
Ag-Pd
28
BaCe0.85Gd0.15O3-a (BCGO)
Ag-Pd
29,30
BaCe0.85Y0.15O3-a (BCY)
Ag-Pd
31
BaCe0.85Dy0.15O3-a (BCD)
Ag-Pd
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BaCe0.9Ca0.1O3-a (BCC)
Ag-Pd
33
La1.9Ca0.1Zr2O6.95 (LCZ)
Ag-Pd
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La1.95Ca0.05Ce2O7-δ (LCC)
Ag-Pd
35
La1.95Ca0.05Zr2O7-δ (LCZO)
Ag-Pd
35
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
Ce0.8La0.2O2-δ (LDC)
Ag-Pd
36
Ce0.8Gd0.2O2-δ (GDC)
Ag-Pd
36
Ce0.8Sm0.2O2-δ (SDC)
Ag-Pd
36
Ce0.8Y0.2O2-δ (YDC)
Ag-Pd
36
(Ce0.8La0.2)0.975Ca0.025O3-δ (CLC)
Ag-Pd
37
YDC-[Ca3(PO4)2 -K3PO4]
(YDCPK)
Ag-Pd
38
BaxCe0.8Y0.2O3 -a + 0.04ZnO
(BCYZ)
Ag-Pd
39
BaCe0.9-xZrxSm0.10O3-α (BCZS)
Ag-Pd
40
BaCe0.8Gd0.1Sm0.1O3-a (BCGS)
Ag-Pd
25
CoFe2O4 (CFO)-Ag
41
Ba0.5Sr0.5Co0.8Fe0.2O3- δ (BSCF)
42
La0.6Sr0.4Fe0.8Cu0.2O3-δ-SDC
43
(Li-Na-K)Carbonate-LiAlO2
BaCe0.85Y0.15O3-a (BCY)
SDC-(Li/Na/K)2CO3
(LSFCu-SDC)
Nafion
SmFe0.7Cu0.1Ni0.2O3 (SFCN)
44
Nafion
Sm1.5Sr0.5NiO4 (SSN)
45
Nafion
SmBaCuFeO5-δ (SBCF)
46
Nafion
SmBaCuCoO5-δ (SBCC)
46
Nafion
SmBaCuNiO5-δ (SBCN)
46
Nafion
Pt/C
47
Sm1.5Sr0.5NiO4 (SSN)
45,48
Sm0.5Sr0.5CoO3-δ (SSCO)
45,48
Sulfonated Polysulfone (SPSF)
SPSF
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3. Progress with H+ Conducting Cells
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In the industrial production of ammonia, a significant fraction of the cost lies in the
preparation and purification of the feed gas [2, 5]. The catalyst is easily poisoned by oxygen and
sulphur compounds. Hydrogen is usually produced from steam reforming of natural gas, which
contains such compounds. Therefore, alternative hydrogen sources, which could eliminate natural
gas, would reduce the overall cost of ammonia and the carbon footprint [2, 16].
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
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Using Nafion as a proton conductor and a Ru/C cathode, Kordali et al. [12] produced NH3
from H2O and N2 at 90oC. The same approach was successfully tested at elevated temperatures
using a perovskite type of proton conductor; the reactants were steam and gaseous nitrogen [49].
In 2007, Wang et al. [50] synthesized ammonia using natural gases (CH4 or C2H6) rather than H2
as a hydrogen source. Recently, Lan et al., reported NH3 synthesis from air and H2O at room
temperature [47, 51].
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Ammonia is a pure hydrogen source for fuel cell applications. To this end, solid state H+
conducting cells were used to study the decomposition of NH3 on Ru [52], Fe [53] and Ag [54]
electrodes. Moreover, the idea to use NH3 as a fuel was tested in both H+ and O2- solid electrolyte
cells with promising results [55, 56].
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The above progress with solid state proton conductors is summarized in Table 2. The solid
electrolyte and the working electrode used in each study are shown in the first and second column,
respectively. The third column shows the reactants and products and the last column contains the
reference(s) of each work.
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Table 2: The reactants and products formed in electrochemical ammonia synthesis and ammonia
utilization cells using alternative hydrogen sources
Solid Electrolyte
Working Electrode
Reactants & Products
Reference
Nafion
Ru/C
N2, H2O ==> NH3 , O2
12
SZY
Ru
N2, H2O ==> NH3 , O2
49
mixed NH4+/H+ Nafion
Pt
N2 (Air), H2O ==> NH3, O2
47
mixed H+/Li+/NH4+
Pt/C
N2 (Air), H2O ==> NH3, O2
51
N2, CH4, C2H6 ==> NH3
50
YDCPK
Ag-Pd
SZY
Ru
NH3 ==> N2, H2
52
CZI
Fe
NH3 ==> N2, H2
53
SCY
Ag
NH3 ==> N2, H2
54
BCG, BCGP
Pt
NH3, O2 ==> N2, H2O
55, 56
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4. Electrochemical Synthesis of Ammonia in Other (not H+) Electrolyte Cells
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As already mentioned, plants can convert molecular N2 into NH3 with a process that
involves metalloenzymes, protons and electrons. Hence, this natural process should also be
considered an “electrochemical” synthesis. Similarly, laboratory studies of NH3 synthesis using
transition metal complexes [8-10] or organic solvent based electrolytes [57] should be considered
as electrochemical processes as well. All these works, however, are not discussed here. The present
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
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review is limited to “classical” electrochemical processes, in which an electrochemical cell is used
and NH3 is formed either from H2 and N2 or from H2O and N2.
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The use of Η2O, instead of Η2, inspired an alternative approach to NH3 synthesis, which is
based on the use of oxygen ion (O2-) conductors. In this case, the solid electrolyte cell is
schematically shown in Figure 2. Steam and nitrogen are introduced together at the cathode, where
the reactions that take place can be summarized as:
3 H2O + 6 e-
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and
3 H2 + N2
158
At the anode, O2- is converted to O2:
<==> 3 O2- + 3 H2
<==> 2 ΝΗ3
3 O2- <==> 3/2 O2 + 6 e-
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(4),
(5).
(6)
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Figure 2: Schematic Diagram of NH3 Synthesis in an O2- Cell.
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and thus, the overall reaction is:
3 H2O + N2 <==> 2 ΝΗ3 + 3/2 O2
164
(7).
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The feasibility of a process in which ΝΗ3 is electrochemically produced from Η2O (steam)
and N2 in O2- conducting solid electrolyte cells, has been successfully demonstrated [49]. The rate
of ΝΗ3 formation was negligibly small at 500oC, but at higher temperatures, an increase by up to
two orders of magnitude was observed.
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Although the reaction rates observed were lower than those obtained in H+ cells, the use of
O conductors offers an additional advantage: pure O2, already separated from the reacting
mixture, is the by-product.
2-
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In the last decade, a number of researchers studied the electrochemical synthesis of NH3 in
molten salt systems [59-63]. In such a molten salt system, the reactions on the two electrodes can
be written as:
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Cathode:
176
Anode:
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so that (1) is the overall reaction. Recently [62], NH3 synthesis was studied in a molten LiCl-KClCsCl system at temperatures between 360-390oC. It was found that a portion of NH3 was
chemically dissolved in the melt in the form of imide (NH2−) and amide (NH2−) anions resulting in
a decrease of the NH3 yield. Despite side reactions, conversions of Li3N into NH3 as high as 70%
were obtained.
<==> N3-
(8),
<==> ΝΗ3 + 3 e-
(9),
1/2 N2 + 3 eN3- + 3/2 H2
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5. Electrochemical Promotion during SSAS
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Solid electrolytes can be used for the electrochemical promotion of the rates of heterogeneous
catalytic reactions. This can be achieved if the catalyst to be promoted is the working electrode of
the solid electrolyte cell. For the reaction of ammonia synthesis, the phenomenon of
electrochemical promotion can be studied in the H+ cell of Figure 1. If the circuit is open (zero
current) and a gaseous mixture of N2 and H2 is introduced over the cathode, the cell will operate as
a catalytic reactor. Thus, ammonia will be formed under open circuit, at a rate denoted as ro (and
expressed in moles of hydrogen reacting per second). When the circuit is closed and a current I is
imposed, the cell will additionally function as an electrochemical proton “pump”. The reaction rate
will increase from ro to r. The effect of electrochemical “pumping” can be quantified by the use of
two dimensionless parameters [63, 64]:
194
(10)
195
and
(11),
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197
with Λ being the Faradaic efficiency, ρ the rate enhancement ratio and Δr the difference between
.
closed and open circuit rate:
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If Λ = 1, the effect is Faradaic, i.e. the increase in reaction rate equals the rate of H+ transport
through the electrolyte. If Λ > 1, the effect is non-Faradaic. In the past 30 years, the phenomenon of
Non-Faradaic Electrochemical Modification of Catalytic Activity (NEMCA), also called
Electrochemical Promotion of Catalysis (EPOC), has been observed in many catalytic reaction
systems [63-65] including oxidations, reductions, hydrogenations, decompositions, and
isomerizations. Values of Λ and ρ as high as 3x105 and 1400, respectively, have been reported [64,
65]. Most of the experimental and theoretical NEMCA studies conclude that this phenomenon is due
to the introduction of charged species from the solid electrolyte to the catalyst/gas interface, followed
by the formation of a neutral double layer. This double layer alters the work function of the catalyst
surface resulting in a reversible alteration of the catalytic reaction rate [63-65]. NEMCA is a powerful
tool in heterogeneous catalysis because a) the surface modification can be monitored
electrochemically and in a reversible manner and b) the increase in the reaction rate can exceed the
rate of ion supply by several orders of magnitude [11, 65].
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
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The phenomenon of Electrochemical Promotion has been investigated in only a few SSAS
works. Marnellos et al. studied the reaction on a Pd electrode and observed very low Λ and ρ
values, i.e. Λ < 2.0 and ρ < 1.5 [15, 21]. Yokari et al. used an industrial Fe catalyst and measured
ρ values as high as 13 when protons were “pumped” towards the catalyst surface. On the contrary,
upon “pumping” H+ away from the catalyst, it was possible to obtain a complete loss of the catalytic
activity [22]. Pitselis et al. studied the reverse reaction (ammonia decomposition) on a Fe catalyst.
A significant decrease in the reaction rate was observed upon pumping H+ to the catalyst surface.
Values of Λ and ρ up to 150 and 3.5 respectively, were observed [53]. Skodra and Stoukides, who
studied the decomposition of ΝΗ3 using Ru as working electrode, observed Λ’s as high as 4.0 when
pumping H+ away from the catalyst surface [52]. Zisekas et al. studied the reaction on Ag
electrodes. Although the Λ’s remained near unity in all experiments, ρ values as high as 57 were
achieved [54].
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A general observation is that, as opposed to catalytic oxidations, hydrogenation reactions, such
as ammonia synthesis, exhibit weak NEMCA characteristics [63-66]. Recently, Garagounis et al.
[66] presented a thermodynamic analysis of NEMCA for reactions with limited equilibrium
conversion. The reaction of ammonia synthesis was used as a model system. The model predictions,
which were in agreement with the existing experimental results, indicated that for ammonia
synthesis at temperatures between 500-600oC, Λ could not exceed the value of 10. The physical
explanation of the model predictions is that, as opposed to reactions with highly negative free
energy change, the effect of H+ pumping on equilibrium limited reactions is both, electrochemical
and catalytic. The pumped protons not only modify the catalytic properties of the WE, but also
serve as the carriers of the electrical power required for ammonia synthesis [66] thus reducing their
promotional effect.
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6. Comparing Catalytic and Electrocatalytic Synthesis
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The catalytic reaction of ammonia synthesis from its elements has been studied extensively in
the past hundred years. Today, the mechanism of the catalytic synthesis is well understood [8] and
consequently, a considerable improvement and optimization of the Haber-Bosch process has been
achieved. On the other hand, the electrocatalytic synthesis has been investigated in the last fifteen
years and quite a few questions need to be answered.
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A comparison of the two processes could start by listing the factors that affect the rate of
ammonia formation in each case. In addition to the reactant feed composition, the inlet flowrate
and the catalyst used, the catalytic rate depends on the pressure and the temperature of operation.
Under the conditions employed in the industrial practice, the reaction conversion is
thermodynamically limited. An increase in temperature increases the reaction rate but, at the same
time, decreases the equilibrium conversion. An increase in pressure increases the equilibrium
conversion. The catalyst accelerates both, the rate of formation and the rate of decomposition and
thus, the composition of the exit stream is close to that predicted by thermodynamics.
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In the electrocatalytic synthesis, the effect of temperature is similar to that in the conventional
catalytic process. The effect of pressure, however, is quite different. When the reaction is carried
out in the cell-reactor of Fig. 1, the half cell reaction that takes place on the working electrode
(catalyst) is reaction (3). In other words, reaction (1), in which 4 volumes of reactants produce 2
volumes of product, is replaced by reaction (3), in which 1 volume of reactant produces 2 volumes
of product. Hence, the effect of pressure is essentially reversed [17, 18].
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
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In addition to the above, several additional factors affect the electrocatalytic reaction rate. The
effects of these additional factors have been presented in the review by Amar et al. [2] and are
briefly summarized here:
258
a) Effect of cell potential.
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There is a minimum voltage difference, below which it is not possible to produce NH3
electrochemically. This threshold value can be calculated from thermodynamics. If ΔGRo is the
Gibbs free energy change for reaction (1), the corresponding standard emf of the cell reaction, Eo,
will be given from the equation [67]:
263
(12)
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271
where F is Faraday’s constant and n is the number of electrons passed per mole of N2 reacted.
According to equations (2) and (3), n = 6. Using equation 12, the values of Eo at 25oC and 500oC
are found to be 0.057 V and -0.123 V, respectively. If the hydrogen source is H2O rather than H2
(reaction 7), the standard corresponding emf values are -1.13 V and -1.21 V at 25oC and 500oC,
respectively. The effect of imposed potential on the rate of formation of NH3 was studied in only a
few of the works of Table 1 [2]. According to Li et al. [30], who studied the reaction at 480oC, the
reaction rate increased by increasing the potential from 0 to 0.6 V while at higher voltages, upon
increasing the potential, the rate gradually leveled off [2, 30].
272
b) Effect of current.
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The effect of applied current is similar to that of the potential. The reaction rate increases with
the imposed current up to a certain value above which, the rate is essentially current independent
[2, 31]. This is because at higher currents the coverage of H+ on the catalyst surface, or the three
phase boundary, is much higher and thus protons combine with each other rather than with Nads,
forming H2 rather than NH3. Another reason could be that the competition between H and N species
on the catalyst surface. In fact, the effects of current and voltage are coupled because, in a cell
operating at certain temperature and reactant composition, it is not possible to change the current
without affecting the cell potential and vice versa.
281
c) Effect of electrolyte.
282
283
284
285
286
287
288
As opposed to the catalytic process, in the electrocatalytic synthesis, one of the reactants is
supplied electrochemically. Thus, the electrolyte properties and in particular, the protonic
conductivity, are crucial because it determines the maximum rate of proton supply. The
conductivity of a solid electrolyte increases exponentially with temperature and is inversely
proportional to the thickness of the electrolyte [68]. By reducing the electrolyte thickness, the same
proton flux is attainable at a lower temperature. Reducing the thickness, on the other hand, may
reduce mechanical stability.
289
d) Effect of working electrode (catalyst).
290
291
292
293
The working electrode must meet the requirements of both, catalytic activity and electric
conductivity. As electrode material, it should exhibit acceptably high electronic conductivity. The
conductivity of the industrial NH3 synthesis catalysts is very poor and this is a problem to be solved
before application in an SSAS system.
294
295
Table 1 shows that SSAS has been studied in a large number of solid electrolyte cells. More
than thirty H+ conducting materials have been tested including perovskite oxides, pyrochlores and
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
296
297
298
299
300
301
302
303
304
305
fluorites [2, 16]. Strictly speaking, polymer electrolytes cannot be considered solid state materials
[2]. Polymer electrolytes such as Nafion and SPSF, however, have been included in Table 1 for
comparison purposes. Table 1 also shows that at least fifteen different materials were used as
working electrodes (WE). The first catalyst tested was palladium [15] because it was assumed that
the cathode should also act as an effective hydrogen "trap", i.e., the electrochemically supplied
hydrogen would not desorb to form gaseous H2 but react on the surface to form NH3 [17, 18].
Following Pd, several catalysts have been tested in the last fifteen years, including Ru, Fe, Pt, AgPd, as well as conductive oxides and composite materials such as SSN, BSCF, SSCO, etc. The AgPd cathode was used in almost half of the SSAS studies, although most of the low temperature
work used a Sm based mixed oxide (Table 1).
306
307
308
309
310
311
312
313
314
315
Table 3 contains information on both, rate of NH3 production and electrolyte conductivity for
studies presented in Tables 1 and 2. The first and second column show the reference and operating
temperature of each work. The third and fourth column contain the type of working electrode and
electrolyte used in each study. The fifth column shows the electrical conductivity of the electrolyte
(in S/cm), calculated at the operating temperature. The sixth column shows the catalytic activity of
the working electrode. The reaction rate, rNH3, is expressed in moles of NH3 produced per second and
per cm2 of electrode area. The last column shows the % Faradaic Efficiency, FE (%), which is defined
as the percentage of protons that is converted to ammonia. To calculate FE (%), one has to divide
rNH3 by the maximum rate of NH3 formation, rmax, for a given proton flux. The latter can be calculated
from the equation:
316
317
318
319
∙
∙
∙ ∙
∙ ∙
(13),
where A is the surface area of the working electrode (in cm2) , I is the current and tH+ is the proton
(H+) transport number, i.e. the fraction of the total charge that is carried by protons. Using equation
13, the FE (%) will be:
%
320
∙ ∙ ∙
∙
∙ 100
(14).
321
322
Values for tH+ were reported in only a few of the SSAS studies of Table 1. Where not reported, it
was assumed that tH+ = 1.
323
324
325
326
327
328
329
330
331
332
There are certain points that are worth mentioning. The first point refers to the catalytic activity
of Pd, Ag-Pd and Pt. The industrial ammonia synthesis catalysts are Fe and Ru based materials.
Some bimetallic systems have also been studied, with a combination of Fe and Co exhibiting the
best results [1]. Table 3, however, shows that at temperatures between 450 and 600oC, the Pd-Ag
electrode exhibits high electrocatalytic activity as well as substantially high Faradaic Efficiency
(50-80 %). One possible explanation could be that the mechanism of the electrocatalytic synthesis
is different from that of the catalytic. In the Haber-Bosch synthesis, it is well established that
nitrogen and hydrogen do not react until the strong triple bond of the N2 molecule has been broken
(dissociative mechanism). In the low temperature metalloenzyme synthesis, however, nitrogen
molecules are hydrogenated without prior rupture of the N-N bond (associated mechanism) [7, 8].
333
Table 3: Reaction Rates and Faradaic Efficiencies
Reference
T (oC)
Working
Electrode
Electrolyte
Conductivity
(S/cm)
rNH3 (mol/s.cm2)
FE ( % )
38
650
Ag-Pd
YDCPK
2.3*10-2
9.5*10-9
≈ 40
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
15, 21
570
Pd
SCY
1.9*10-3
4.5*10-9
78
26
550
Ag-Pd
LSGM
3.9*10-3
2.37*10-9
≈ 70
32
530
Ag-Pd
BCD
0.93*10-2
3.5*10-9
≈ 80
42
530
BSCF
BCY
7*10-3
4.1*10-9
60
28
520
Ag-Pd
LBGM
0.82*10-2
1.89*10-9
≈ 60
31
500
Ag-Pd
BCY
5.4*10-3
2.1*10-9
≈ 60
40
500
Ag-Pd
BCZS
4*10-3
2.67*10-9
≈ 80
33
480
Ag-Pd
BCC
2.34*10-4
2.69*10-9
≈ 50
43
450
LSFCu-SDC
SDCcarbonate
3.3*10-3
5.39*10-9
≈ 3.5
41
400
CFO-Ag
CarbonateLiAlO2
4*10-5
2.32*10-10
≈ 2.5
12
90
Ru/C
Nafion
4.5*10-2
2.12*10-11
<1
45
80
SSN
Nafion
4.5*10-2
1.05*10-8
<1
45
80
SSN
SPSF
5*10-2
1.03*10-8
<1
44
80
SFCN
SPSF
5*10-2
1.13*10-8
90.4
46
80
SBCN
Nafion
4.5*10-2
8.7*10-9
<1
47
25
Pt
mixed
NH4+/H+
Nafion
5*10-2
3.1*10-9
(N2+H2)
2.2
47
25
Pt
mixed
NH4+/H+
Nafion
5*10-2
1.14*10-9
<1
H+/Li+/NH4+
Mixed
conducting
membrane
5.1*10-3
51
80
Pt-C
(Air + H2O)
9.37*10-10
<1
334
335
336
337
338
339
340
The experimental information obtained so far on SSAS is not sufficient to favor either the
associative or the dissociative mechanism. Skulason et al, however, published recently a theoretical
evaluation of metal electro-catalysts for electrochemical ammonia synthesis. The most active
surfaces were Fe, Ru, Mo and Rh, although the formation of gaseous H2 reduced the Faradaic
Efficiency. The catalytic activity of platinum and palladium were among the lowest for this
reaction. It was also found that early transition metals such as Sc, Y, Ti and Zr bind N-adatoms
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
341
342
343
344
345
346
347
348
349
stronger than H-adatoms and therefore, upon imposing a negative voltage, these metals can
effectively catalyze ammonia synthesis [8]. The catalytic activity of early transition metals may
explain the high reaction rates and Faradaic Efficiencies reported for Pd and Ag-Pd. Most of the
solid electrolytes used in the works of Table 1 (e.g. SZY, LSGM, SCY, BCY) contain transition
metals such as Y, Zr, Ce, Yb, etc. Moreover, under reaction conditions (closed circuit, 450-600oC),
several changes may occur in the structure of these materials. For example, in the SrCe0.95Yb0.05O3-a
(SCY) electrolyte used in reference 15, EDS analysis showed that Sr and Yb had migrated from
the bulk of the electrolyte to the Pd-SCY interphase. Thus, the observed activity of Pd and Ag-Pd
could be attributed to the presence of early transition metals at the electrode-electrolyte interphase.
350
351
352
353
354
355
356
357
358
359
The second point refers to the dependence of rNH3 and FE on temperature. The highest rates of
NH3 formation were obtained at low temperatures (25-90oC), while, with the exception of reference
44, at temperatures between 500 and 650oC, the FE values are at least one order of magnitude
higher than those attained with Nafion and SPSF, at low temperatures (25-90oC). The conductivity
values (fifth column of Table 3) can give an explanation for this. The conductivity of Nafion and
SPSF is much higher than that of the ceramic high temperature conductors. On the other hand, the
catalytic reaction rate increases with temperature. Thus, at low temperatures, only a small fraction
of the electrochemically supplied protons react to produce ammonia. At high temperatures, the
catalytic activity increases and thus, FE increases as well. The rates of H+ transport through the
ceramic electrolytes, however, remain much lower than those in Nafion and SPSF.
360
361
7. Economic Considerations and Future Outlook
362
363
364
365
366
367
368
369
370
Ammonia is one of the top-ranking chemicals in terms of volume, primarily due to its use in
the fertilizer industry. Its global annual production is well in excess of 100 million tons with a
tendency to increase further [2]. It is also, a significant component of the world energy budget, as
its production accounts for more than 1% of man-made power consumption. Furthermore, as the
combustion of ammonia as opposed to carbon-containing fuels, releases only H2O and N2 to the
environment, it has been proposed as a practical alternative to fossil fuels for internal combustion
engines and even fuel cells, either directly or after decomposition [69]. Also, since ammonia is an
excellent carbon-free hydrogen storage medium, SSAS can be combined with plants which produce
energy from renewable sources (solar or wind energy).
371
372
373
374
375
376
377
378
379
The interest in all methods of ammonia synthesis, including SSAS, is thus expected to continue
to increase. There are certain advantages of SSAS vs. the conventional Haber-Bosch process. A
large percentage of the cost in the industrial production of ammonia lies in the preparation and
purification of the reactant gases. Water vapour and carbon monoxide can cause a reversible
poisoning if their concentrations are small, while trace amounts of oxygen and most sulphur
compounds cause irreversible poisoning of the catalyst [3-5]. Therefore, extensive and costly prepurification of the nitrogen (O2 removal) and hydrogen (H2O and CO removal) is necessary to
minimize poisoning. In SSAS, with hydrogen source being either H2O or H2, only one of the
reactants, N2, needs purification.
380
381
382
383
384
385
386
On the other hand, several technical and material requirements must be met in SSAS [2, 11,
70]. Firstly, the catalyst will also serve as the working electrode and therefore must exhibit, not
only high catalytic activity towards ammonia synthesis, but also adequately high electronic
conductivity. Precious metals based on Pd and Ru, have been found to satisfy these demands. If
these metals are to be used, then the quantity of catalyst required per unit cell should be kept to a
minimum [16]. Secondly, in addition to mechanical strength and chemical stability, the electrolyte
material should exhibit high protonic conductivity at the temperature of operation. As seen in Table
Garagounisetal.
ElectrochemicalSynthesisofAmmonia
387
388
389
390
391
3, there are no materials with high protonic conductivity at temperatures between 100oC and 500oC.
This is unfortunate because at low temperatures, although the protonic conductivity of certain
materials is substantially high, low catalytic reaction rates are obtained, while at high temperatures
the Faradaic Efficiency can be acceptably high, but the reaction rate is limited by the proton
conductivity of the electrolyte.
392
393
394
395
396
397
398
399
400
401
402
In the past 15 years, many researchers worldwide have contributed to the improvement of the
SSAS process. New proton conductors have been prepared and tested and catalytically effective
working electrodes have been employed. Notably, Lan et al. [47] recently synthesized ammonia
from air and water at ambient temperature using Pt electrodes. Robinson et al. [71, 72] fabricated
high temperature (650-850oC), thin, tubular ceramic proton conductors through which they passed
as many as 5.5×10-6 mol H2/s.cm2, while Xu et al. [44] reported the highest ammonia production
rate so far, i.e. 1.13 x 10-8 mol·cm-2·s-1, when employing a Nafion membrane and a SmFe0.7Cu0.3xNixO3 cathode. Along the lines of electrolytes with high conductivity at intermediate temperatures
(150-300 oC) the recent work of Martsinkevich and Ponomareva [73], who fabricated electrolyte
materials exhibiting maximum proton conductivity at 200-220oC, offers an alternative that is worth
investigation.
403
404
405
406
407
408
409
If current densities similar to those achieved in fuel or electrolysis cells (0.25 - 0.5 A·cm-2)
were achievable simultaneously with a FE(%) of approximately 50%, then commercially viable
production rates (in the range 4.3 - 8.7 x 10-7 mol·cm-2·s-1) could be attained [16]. The lack of
information concerning long term stability of most materials and electrochemical cells in ammonia
containing environments is partly due to the relatively early stage of development of this
technology [16]. Such information, along with the actual cost of fabrication, would be necessary
for a comprehensive economic assessment (life-cycle analysis) of a SSAS plant.
410
411
Acknowledgements
412
413
414
We gratefully acknowledge financial support of this research by the European Union and the
General Secretariat of Research and Technology of Greece (Program ARISTEIA, Project Number
1089).
415
416
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